2- Oxidation reduction (Redox) titration

Redox titrations are based on a reduction-oxidation reaction between an

oxidizing agent and a reducing agent. A potentiometer or a redox indicator is

usually used to determine the endpoint of the titration, as when one of the

constituents is the oxidizing agent potassium dichromate. The color change of

the solution from orange to green is not definite, therefore an indicator such as

sodium diphenylamine is used. Analysis of wines for sulfur dioxide requires

iodine as an oxidizing agent. In this case, starch is used as an indicator; a blue

starch-iodine complex is formed in the presence of excess iodine, signalling the

endpoint.

Some redox titrations do not require an indicator, due to the intense color

of the constituents. For instance, in permanganometry a slight persisting pink

color signals the endpoint of the titration because of the color of the excess

oxidizing agent potassium permanganate.[31] In iodometry, at sufficiently large

concentrations, the disappearance of the deep red-brown triiodide ion can itself

be used as an endpoint, though at lower concentrations sensitivity is improved

by adding starch indicator, which forms an intensely blue complex with

triiodide.

Assay

An assay is a form of biological titration used to determine the

concentration of a virus or bacterium. Serial dilutions are performed on a sample

in a fixed ratio (such as 1:1, 1:2, 1:4, 1:8, etc.) until the last dilution does not

give a positive test for the presence of the virus. The positive or negative value

may be determined by visually inspecting the infected cells under

a microscope or by an immunoenzymetric method such as enzyme-linked

immunosorbent assay (ELISA). This value is known as the titer.

Oxidation-Reduction (Redox)

3 Fe2+ + Cl2 = 2Fe3+ + 2Cl

3Fe2+ - 2e = 2Fe3+ (Loss of electrons: Oxidation)

Cl2 + 2e = 2Cl- (Gain of electrons: Reduction)

In every redox reaction, both reduction and oxidation must occur.

Substance that gives electrons is the reducing agent or reductant. Substance that

accepts electrons is the oxidizing agent or oxidant.

Reaction of ferrous ion with chlorine gas Overall, the number of

electrons lost in the oxidation half reaction must equal the number gained in the

reduction half equation.

Oxidation Number (O.N)

The O.N of a monatomic ion = its electrical charge.

The O.N of atoms in free un-combined elements = zero

The O.N of an element in a compound may be calculated by assigning the O.N

to the remaining elements of the compound using the aforementioned basis and

the following additional rules:

The O.N. for oxygen = -2 (in peroxides = -1).

The O.N. for hydrogen = +1 (in special cases = -1).

The algebraic sum of the positive and negative O.N. of the atoms represented by

the formula for the substance = zero.

The algebraic sum of the positive and negative O.N. of the atoms in a

polyatomic ion = the charge of the ion.

Balancing Redox Reactions using Half Reaction Method

1- Divide the equation into an oxidation half-reaction and a reduction

half-reaction

2- Balance these – Balance the elements other than H and O

3- Balance the O by adding H2O – Balance the H by adding H+

4- Balance the charge by adding e-

5- Multiply each half-reaction by an integer such that the number of

electron lost in one equals the number gained in the other

6- Combine the half-reactions and cancel

7- Balancing Redox Reactions Using the Half Reaction Method

Example-1

Al + HCl = AlCl3 + H2

0 +1 -1 +3 -1 0

Oxidation: (Al = Al3+ + 3e-) 2 = 6 e–

Reduction: (2H+ + 2e = H2 ) 3 = 6 e–

Redox: 2Al + 6H+ = 2Al3+ + 3H2

Example-2

MnO4 + C2O4

2 + H+ = Mn2+ + CO2 + H2O

1- Balance each half reaction:

MnO4 + 5é = Mn2+

C2O4 2 = 2 CO2 + 2é

2- Use the number of moles so as to make the electrons gained in one

reaction equal those lost in the other one

2MnO4 + 5C2O4

2 = 2 Mn2+ + 10CO2

3- Balance oxygen atoms by adding water

2MnO4 + 5C2O4

2 = 2Mn2+ + 10CO2 + 8H2O

4- Balance hydrogen atoms by adding H+

2MnO4 + 5C2O4

2 + 16H+ = 2Mn2+ + 10CO2 + 8H2O

Nernest Equation for Electrode Potential

Et = Eo + RT/nF log [Mn+]

E25 °C = Eo + 0.0591/n log [Mn+]

Where

Et = electrode potential at temperature t.

E = standard electrode potential (constant depend on the system)

R = gas constant

T = absolute Temp. (t °C + 273)

F = Faraday (96500 Coulombs) loge = ln (natural logarithm= 2.303 log)

n= valency of the ion

[Mn+] = molar concentration of metal ions in solution

Standard Electrode Potential (Eo)

Eo is the electromotive force (emf) produced when a half cell

(consisting of the elements immersed in a molar solution of its ions) is coupled

with a standard hydrogen electrode (E = zero).

Li / Li+ –3.03 Cd/Cd2+ –0.40 K / K+ –2.92

Sn/Sn2+ –0.13 Mg/Mg2+ –2.37 H2 (Pt)/H+ 0.00

Al/Al3+ –1.33 Cu/Cu2+ +0.34 Zn/Zn2+ –0.76

Hg/Hg2+ +0.79 Fe/Fe2+ –0.44 Ag/Ag+ +0.8

Detection of End Point in Redox Titrations

1- Self Indicator (No Indicator)

KMnO4 (violet) + Fe2+ + H+ Mn2+ (colourless) + Fe3+

2- External Indicator

Cr2O7 2 + 3Fe2+ + 14H+ 2Cr3+ + 3Fe3+ + 7H2O

3- Internal Redox Indicator

Inox + n e = Inred

4- Irreversible Redox Indicators

Some highly coloured organic compounds that undergo irreversible

oxidation or reduction e.g. Methyl Orange

Properties of Oxidizing Agents

1. Potassium permanganate (KMnO4)

2. Potassium dichromate (K2 Cr2O7 )

3. Iodine (I2)

4. Potassium iodate (KIO3)

5. Bromate-bromide mixture

Measuring the endpoint of a titration

Different methods to determine the endpoint include:

Indicator: A substance that changes color in response to a chemical

change. An acid–base indicator (e.g., phenolphthalein) changes color depending

on the pH. Redox indicatorsare also used. A drop of indicator solution is added

to the titration at the beginning; the endpoint has been reached when the color

changes.

Potentiometer: An instrument that measures the electrode

potential of the solution. These are used for redox titrations; the potential of the

working electrode will suddenly change as the endpoint is reached.

An elementary pH meter that can be used to monitor titration reactions

pH meter: A potentiometer with an electrode whose potential

depends on the amount of H+ ion present in the solution. (This is an example of

an ion-selective electrode.) The pH of the solution is measured throughout the

titration, more accurately than with an indicator; at the endpoint there will be a

sudden change in the measured pH.

Conductivity: A measurement of ions in a solution. Ion

concentration can change significantly in a titration, which changes the

conductivity. (For instance, during an acid–base titration, the H+ and OH− ions

react to form neutral H2O.) As total conductance depends on all ions present in

the solution and not all ions contribute equally (due to mobility and ionic

strength), predicting the change in conductivity is more difficult than measuring

it.

Color change: In some reactions, the solution changes color without

any added indicator. This is often seen in redox titrations when the different

oxidation states of the product and reactant produce different colors.

Precipitation: If a reaction produces a solid, a precipitate will form

during the titration. A classic example is the reaction between Ag+ and Cl− to

form the insoluble salt AgCl. Cloudy precipitates usually make it difficult to

determine the endpoint precisely. To compensate, precipitation titrations often

have to be done as "back" titrations (see below).

Isothermal titration calorimeter: An instrument that measures the

heat produced or consumed by the reaction to determine the endpoint. Used

in biochemical titrations, such as the determination of how substrates bind

to enzymes.

Thermometric titrimetry: Differentiated from calorimetric titrimetry

because the heat of the reaction (as indicated by temperature rise or fall) is not

used to determine the amount of analyte in the sample solution. Instead, the

endpoint is determined by the rate of temperature change.

Spectroscopy: Used to measure the absorption of light by the

solution during titration if the spectrum of the reactant, titrant or product is

known. The concentration of the material can be determined by Beer's Law.

Amperometry: Measures the current produced by the titration

reaction as a result of the oxidation or reduction of the analyte. The endpoint is

detected as a change in the current. This method is most useful when the excess

titrant can be reduced, as in the titration of halides with Ag+.

1- Redox reactions are those reactions which involves change in the

oxidation number or transfer of electrons between reactants.

2- Oxidation is a process which involves loss of electrons.

Fe2+ - e = Fe3+

It is a process in which an atom or an ion may lose one or more electrons.

3- Reduction is a process in which electrons are gained.

Cl2 + 2e = 2Cl-

It is a process in which an atom or an ion may gain one or more electrons.

4- The equivalent weight of an oxidizing agent (oxidant) or reducing

agent (reductant) is equal to its molecular weight divided by the number

of electrons lost or gained by this substance during the reaction:

MnO4− +8H+ +5e = Mn2+ +4H2O

The eq.wt. = MnO4

−

5 =

KMnO4

5

Cr2O7−− + 6e + 14H+ = 2Cr3+ +7H2O

The eq.wt. = Cr2O7

−−

6 =

K2Cr7O7

6

5- Indicators in oxidation - reduction reaction:

Potassium permanganate may act as a self-indicator as it has a distinct

purple color and becomes colorless on reduction. In the case of potassium

dichromate an indicator must be used for the determination of the end point

which may be an internal or an external indicator. The types of indicators used

may be as follows: Self-indicator - Internal indicator - External Indicator

Oxalic acid is oxidized by potassium permanganate , in acid solution to

carbon dioxide and water.

2KMnO4 +3H2S04 + 5H2C2O4 = 2MnSO4 + K2SO4 + 10CO2 + 8H2O

The reaction is complete at a temperature of about 60-90°C

𝐶2𝑂4−2 → 2CO2 + 2e-

Oxidation with potassium permanganate:

KMnO4 is a strong oxidizing agent in acid medium

2KMnO4 + 3H2SO4 = K2SO4+ 2MnSO4 + 3H2O+5O

HCl, could not be used instead of H2SO4 as it is readily oxidized to chlorine in

presence of permanganate.

2KMnO4 + 16HCI = 2KCl + 2MnCl2 + 5Cl2 +8H2O

Nitric acid is stronger than KMnO4

In strong alkaline medium, heptavalent manganese is reduced as follows:

2KMnO4 = K2O + 2MnO2 + 3O

And the eq.wt. in alkaline sol.=1/3 the molecular weight. The formed

manganese dioxide which is black in color may mask the end point in alkaline

medium. KMnO4, however is not a primary standard as it is difficult to obtain in

a purified state due to the fact that it is always contaminated with manganese

dioxide.

4Mn𝑂4− + 2H2O → 4MnO2 + 3O2 + 4OH-

(I) Standardization of sodium thiosulphate solution with potassium iodate:

Theory

Potassium iodide reacts in acid medium with potassium iodate with the

liberation of Iodine:

KIO3 + 3H2SO4 + 5KI → 3K2SO4 + 3I2 ↑ + 3H2O

2Na2S2O3 + I2 → Na2S4O6 + 2NaI

(II) Determination of copper in copper sulphate (CuSO4.5H2O)

Theory

Copper sulphate reacts with potassium iodide according to the following

equation:

2CuSO4 + 4KI → Cu2I2 + I2↑ + 2K2SO4

2Na2S2O3 + I2 → Na2S4O6 + 2NaI

i.e. 2CuSO4 ≡ I2 ≡ 2Na2S2O3

Titration of Hydrogen Peroxide

EXPERIMENTAL GOALS:

The purpose of this lab is to learn how to use a titration to standardize a

solution and to determine the concentration of an unknown solution.

INTRODUCTION:

A chemist or biochemist will often find it necessary to determine the

concentration of something in a solution or sample. One method that can

accomplish this is the method of titration. In a titration, a solution of known

concentration is allowed to react with a solution of unknown concentration until

all of the “known” is consumed in the reaction. By carefully measuring the

volume of each solution, one can easily calculate the concentration of the

unknown solution.

Titration is performed on the basis of stoichiometry. For example,

consider two substances that react on a 2:1 stoichiometric basis. If it is known

that one mole of a reactant reacts on a 2:1 basis with another reactant of known

concentration, then we can use the stoichiometric relationship between the two

reactants to figure out the concentration of the reactant of unknown

concentration. The important consideration here is that the titration must be

stopped at the point where the substances have reacted in stoichiometrically

equivalent amounts. This point is referred to as the equivalence point.

Underestimating or overestimating the volume of either the reactant of known or

unknown concentration during the titration can lead to an inaccurate experiment.

The equivalence point of a titration can be detected by various means. In

today’s lab, we are going to use the reaction between potassium permanganate

and hydrogen peroxide in order to learn more about titrations as well as

stoichiometry. This particular reaction is actually an oxidation/reduction

reaction rather than an acid/base reaction. The sulfuric acid solution (H2SO4) is

being used to provide an acidic environment for the oxidation-reduction reaction

to take place in. Finally, it is important to remember that balancing of some

oxidation/reduction reactions is more complicated than other reactions because

you have to balance mass (# of atoms of each element) as well as charge (related

to the number of electrons lost/gained).

2KMnO4 + 5H2O2 + 3H2SO4 2MnSO4 + K2SO4 + 5O2 + 8H2O

In this particular reaction, we use the faint pink color of the permanganate

ion as an indicator rather than using an additional dye to determine the end

point. In essence, the permanganate in the KMnO4 solution acts as its own

indicator. The Mn2+ ion is actually colorless, but when MnO4- is in excess, the

solution will turn from colorless to pink, indicating that the equivalence point

has been reached.

When performing a titration, it is necessary to first determine the

concentration of the known solution as accurately as possible. This process is

referred to as standardization. For example, in this experiment, the titrant is an

aqueous solution of potassium permanganate, KMnO4; however, solid KMnO4

contains traces of other impurities, such as MnO2, so a solution of known

concentration cannot be made by simply measuring out a certain number of

grams of KMnO4 and dissolving it in the appropriate amount of water. To

determine the concentration of the KMnO4, it is titrated against something that

can be measured accurately, such as solid oxalic acid, H2C2O4, whose mass can

be measured very precisely. By using the stoichiometry of the standardization

reaction, the concentration of the titrant solution can be determined.

2MnO4- + 5H2C2O4 + 6H+ 2Mn2+ + 10CO2(g) + 8H2O

The level of accuracy afforded by graduated cylinders is not sufficient for

a titration, so more accurate instruments must be used. Burets or pipettes

typically are used in titrations. It is important to utilize the significant figures

that a buret allows one to use. Burets are very accurate pieces of equipment. In

order to determine the volume dispensed from a buret, the initial and final

volumes are recorded and subtracted. (See the Introduction for the

“Stoichiometry III: Titration of Vinegar” section for a more thorough discussion

of the use of burets in titrations.)

Determination of strength (g/L) of ferrous ammonium sulphate

(FeSO4.(NH4)2SO4.6H2O) by titrating it against standard (1.0 g/L)

potassium dichromate (K2Cr2O7) solution

K2Cr2O7 + 4H2SO4 → K2SO4 + Cr2(SO4)3 + 4H2O +3O

6 FeSO4 + 3H2SO4 + 3O → 3Fe2(SO4)3 + 3H2O

K2Cr2O7 + 6FeSO4 + 7H2SO4 = 3Fe2(SO4)3 + K2SO4 + Cr2(SO4)3 + 7H2O

(Cr2O7)2- + 6Fe2+ + 14H+ = 6Fe3+ + 2Cr3+

N-phenyl anthranilic acid is used as an indicator. Indicator is not oxidized

as long as Fe2+ ions are there in the solution. The slight excess amount of

dichromate will oxidize the indicator when all of the Fe2+ ions have been

converted to Fe3+ ions resulting in colour change of the solution from greenish

(due to Cr3+) to purple.